Use the periodic table to answer the following questions. (a) Identify an element with five electrons in the outer subshell. (b) Identify an element that would tend to lose two electrons. (c) Identify an element that would tend to gain two electrons. (d) Identify the group having metal, non-metal, liquid as well as gas at the room temperature.

> [!TLDR] > The Aufbau Principle and periodic trends let us pinpoint elements by their outer-subshell electron count, their tendency to lose or gain electrons (metallic/non-metallic character), and the unique Group 17 that contains metals, non-metals, liquids, and gases at room temperature. Let’s unpack each part by first understanding the guiding idea: the **Aufbau Principle** tells us the order in which electrons fill subshells ($1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, \dots$). The outermost subshell (the one with the highest principal quantum number $n$ that contains electrons) determines an element’s chemical behaviour. The number of electrons in that outer subshell — especially for $s$ and $p$ subshells — decides whether the element tends to lose, gain, or share electrons. --- ### (a) An element with five electrons in the outer subshell 1. The outer subshell is the single outermost subshell being filled — for the elements at the right of the table, the $np$ subshell. Five electrons **in that one subshell** means $np^5$, so the configuration ends $ns^2 np^5$. This is the **halogen family** — Group 17. > [!NOTE] > Do not confuse the outer *subshell* with the outer *shell*. Five electrons in the outer SHELL would mean $ns^2np^3$ (Group 15); only $np^5$ puts five electrons in a single outer subshell, which is what the question asks. 2. Examples: Fluorine ($\text{F}$, $1s^2 2s^2 2p^5$), Chlorine ($\text{Cl}$, $[\text{Ne}]\,3s^2 3p^5$), Bromine, Iodine. Any one of these fits. > [!TIP] > The $np^5$ configuration is one electron short of a complete noble-gas shell — which is why the halogens are the table's most eager electron gainers (most negative electron gain enthalpies). **Answer for (a):** Fluorine (or any Group 17 element). --- ### (b) An element that would tend to lose two electrons 1. An element that **loses two electrons** easily is one with two valence electrons in its outer $s$ subshell — i.e., $ns^2$ configuration. These are the **alkaline earth metals*

Use the periodic table to answer the following questions. (a) Identify an element with five electrons in the o

Exercises · 3.27

Q. Use the periodic table to answer the following questions.

(a) Identify an element with five electrons in the outer subshell.

(b) Identify an element that would tend to lose two electrons.

(d) Identify the group having metal, non-metal, liquid as well as gas at the room temperature.

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Concept understanding — Aufbau Principle

The Aufbau Principle: Building Atoms from the Ground Up

Imagine you're building a house. You don't start with the roof — you lay the foundation first, then the walls, then the roof. Electrons in an atom behave the same way: they fill the lowest-energy "rooms" (orbitals) before moving to higher-energy ones. That's the core idea behind the Aufbau Principle.

The word Aufbau comes from German, meaning "building up" or "construction." It describes how electrons are added one by one to an atom as you move across the periodic table.

The Intuition: Why Lowest Energy First

An atom is most stable when its electrons are in the lowest possible energy states. Think of it like a ball rolling down a hill — it naturally settles at the bottom, not halfway up. If you try to put an electron into a high-energy orbital while a lower-energy orbital is empty, the atom would be unstable and the electron would "fall" down to the empty spot, releasing energy.

So nature follows the path of least energy: fill from the bottom up.

The Precise Statement

Aufbau Principle: In the ground state of an atom, electrons fill atomic orbitals in order of increasing orbital energy. Lower-energy orbitals are completely filled before electrons occupy higher-energy orbitals.

This is not a guess — it's a rule that emerges from quantum mechanics and is confirmed by experimental data (like atomic spectra and ionization energies).

The Order of Filling: The (n + l) Rule

But here's where it gets interesting: orbital energies don't simply increase with principal quantum number $n$ . For example, the 4s orbital ( $n=4, l=0$ ) actually has lower energy than the 3d orbital ( $n=3, l=2$ ). So electrons fill 4s before 3d.

To predict the correct order, use the (n + l) rule:

Tip

For an orbital with quantum numbers $n$ and $l$ , its energy increases with $(n + l)$ . If two orbitals have the same $(n + l)$ value, the one with lower $n$ fills first.

Let's check:

4s: $n + l = 4 + 0 = 4$
3d: $n + l = 3 + 2 = 5$

Since $4 < 5$ , 4s fills first. That's why potassium (atomic number 19) has the electron configuration $[\text{Ar}]\,4s^1$ , not $[\text{Ar}]\,3d^1$ .

The Complete Filling Order (Memory Aid)

Here's the sequence you'll use for most elements up to atomic number 118:

1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \rightarrow 4p \rightarrow 5s \rightarrow 4d \rightarrow 5p \rightarrow 6s \rightarrow 4f \rightarrow 5d \rightarrow 6p \rightarrow 7s \rightarrow 5f \rightarrow 6d \rightarrow 7p

A common trick is to write the orbitals in rows by $n$ and draw diagonal arrows:

1s
2s  2p
3s  3p  3d
4s  4p  4d  4f
5s  5p  5d  5f
6s  6p  6d
7s  7p

Follow the diagonals from top-right to bottom-left — that gives you the filling order.

Important Exceptions (Don't Panic)

The Aufbau principle works beautifully for most elements, but there are a few well-known exceptions where a half-filled or fully-filled subshell provides extra stability:

Chromium ( $Z=24$ ): Expected $[\text{Ar}]\,4s^2 3d^4$ , but actual is $[\text{Ar}]\,4s^1 3d^5$ (half-filled d-subshell is more stable) …

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The Aufbau Principle and periodic trends let us pinpoint elements by their outer-subshell electron count, their tendency to lose or gain electrons (metallic/non-metallic character), and the unique Group 17 that contains metals, non-metals, liquids, and gases at room temperature.

Let’s unpack each part by first understanding the guiding idea: the Aufbau Principle tells us the order in which electrons fill subshells ( $1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, \dots$ ). The outermost subshell (the one with the highest principal quantum number $n$ that contains electrons) determines an element’s chemical behaviour. The number of electrons in that outer subshell — especially for $s$ and $p$ subshells — decides whether the element tends to lose, gain, or share electrons.

(a) An element with five electrons in the outer subshell

The outer subshell is the single outermost subshell being filled — for the elements at the right of the table, the $np$ subshell. Five electrons in that one subshell means $np^5$ , so the configuration ends $ns^2 np^5$ . This is the halogen family — Group 17.

Note

Do not confuse the outer subshell with the outer shell. Five electrons in the outer SHELL would mean $ns^2np^3$ (Group 15); only $np^5$ puts five electrons in a single outer subshell, which is what the question asks.

Examples: Fluorine ( $\text{F}$ , $1s^2 2s^2 2p^5$ ), Chlorine ( $\text{Cl}$ , $[\text{Ne}]\,3s^2 3p^5$ ), Bromine, Iodine. Any one of these fits.

Tip

The $np^5$ configuration is one electron short of a complete noble-gas shell — which is why the halogens are the table's most eager electron gainers (most negative electron gain enthalpies).

Answer for (a): Fluorine (or any Group 17 element).

(b) An element that would tend to lose two electrons

An element that loses two electrons easily is one with two valence electrons in its outer $s$ subshell — i.e., $ns^2$ configuration. These are the alkaline earth metals — Group 2.
Losing two electrons gives a stable noble-gas configuration (e.g., $\text{Mg} \rightarrow \text{Mg}^{2+}$ with $1s^2 2s^2 2p^6$ ). The metallic character is high, so they readily form $+2$ cations.
Examples: Beryllium ( $\text{Be}$ ), Magnesium ( $\text{Mg}$ ), Calcium ( $\text{Ca}$ ), etc.

Watch out

Do not confuse with Group 1 (alkali metals) which lose one electron, not two. Group 2 elements lose exactly two because the $ns^2$ pair is the only valence pair.

Answer for (b): Magnesium (or any Group 2 element).

(c) An element that would tend to gain two electrons

An element that gains two electrons wants to complete its octet by adding two electrons to its outer subshell. This happens when the outer subshell has six electrons — $ns^2 np^4$ — so adding two gives $ns^2 np^6$ (noble gas). These are the chalcogens — Group 16.
Examples: Oxygen ( $\text{O}$ , $1s^2 2s^2 2p^4$ ) gains two to become $\text{O}^{2-}$ ; Sulfur ( $\text{S}$ , $3s^2 3p^4$ ) similarly forms $\text{S}^{2-}$ .

Tip

The tendency to gain electrons is strongest for non-metals near the right of the periodic table. Oxygen and sulfur are classic examples; selenium and tellurium also show this, but less strongly.

Answer for (c): Oxygen (or any Group 16 element).

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Q.Use the periodic table to answer the following questions.

(a) An element with five electrons in the outer subshell

(b) An element that would tend to lose two electrons

(c) An element that would tend to gain two electrons

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